Inorganic AS-Level

• How is the modern periodic table arranged?

  The elements are arranged in order of increasing atomic numbers.

  -> this arrangement reveals patterns in the properties of elements, allowing chemists to make predictions about their behaviour.
• How are periods arranged?

  Elements in the same period have the same number of electron shells.
• Why do elements in the same group exhibit similar properties?

  Elements in the same group have similar electron configurations in their outer shell; so they exhibit similar physical and chemical properties.
• What determines which block an element is in?

  Which orbital their outermost electrons are in.

  eg. elements in the s-block have their outermost electrons in an s orbital.
• Where are the blocks in the periodic table?

  
• In period 3, how does melting and boiling point change for the three metals?

  • For the metals (Na, Mg, Al), the melting and boiling points increase across the period due to the increasing strength of the metallic bonding.
• Why does the strength of metallic bonding increase across period 3?

  • The ions have a larger positive charge (1+, 2+, 3+)

  • The ions have a smaller atomic radius.

  • There are ore delocalised electrons (eg. Na only has 1, but Al has 3).

  -> These factors result in stronger electrostatic attraction between metal cations and delocalised electrons, which require greater amounts of energy to overcome.
• Which element in period 3 has the highest melting point? Why?

  Silicon (Si)

  • Giant covalent lattice structure.

  • Each atom is covalently bonded to 4 other atoms in a tetrahedral arrangement.

  • This forms very strong bonds linking all the atoms together in a 3D lattice.

  • A huge amount of energy is required to break these covalent bonds, resulting in a very high melting and boiling point.
• Why do phosphorous (P4), sulfur (S8) and chlorine (Cl2) have low melting and boiling points?

  • These are simple molecular structures.

  • This means there are only weak induced dipole-dipole forces existing between the molecules.

  • These intermolecular forces are easily overcome, requiring little energy, so results in low melting and boiling points.
• Which has the highest melting and boiling point out of the three simple molecular structures in period 3?

  • Sulfur (S8), because its molecules contain more electrons than phosphorous and chlorine.

  • The larger S8 molecules allow stronger induced dipole-dipole forces to form between molecules.
• Which has the lowest melting and boiling point out of the three simple molecular structures in period 3?

  Chlorine (Cl2) has the lowest melting and boiling points among the three because its molecules contain the fewest electrons, resulting in the weakest intermolecular forces between Cl2 molecules.
• Which element in period 3 has the lowest melting and boiling point?

  Argon (Ar) - as it is a noble gas so its atoms do not form any bonds and only very weak induced dipole-dipole forces attract them. Minimal energy is needed to overcome these negligible forces between atoms.
• What is the graph showing melting point across period 3?

  
• What happens to atomic radius across period 3?

  Decreases
• Why does atomic radius decrease across period 3?

  • As protons are added across a period, the nuclear charge increases.

  
  

  • This results in a stronger electrostatic attraction between the nucleus and the outer electrons, drawing the outer electrons closer to the nucleus.

  
  

  • The electrons added across a period go into the same energy level, so shielding is similar.

  
  

  • With increasing nuclear charge and minimal change in shielding, the stronger electrostatic attraction causes the atomic radius to decrease across the period.
• What is ionisation?


  The process of removing one or more electrons from an atom or molecule.
• Is ionisation exothermic or endothermic?

  Endothermic - as ionisation requires an input of energy.

  -> The values for ionisation energies are always positive.
• What is the first ionisation energy?

  The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms, to form 1 mole of gaseous 1+ ions.
• What is the equation for the first ionisation energy of lithium?

  Li(g) -> Li+(g) + e-
• What does the ionisation energy of an atom depend upon?

  How strongly its outermost electrons are attracted to the nucleus (electrostatic attraction).
• What are the three factors affecting ionisation energy (electrostatic attraction)?

  1) Nuclear charge

  2) Atomic radius

  3) Electron shielding
• How does nuclear charge affect ionisation energy?

  Atoms with more protons in their nucleus have a stronger positive charge, which creates a stronger electrostatic attraction between the nucleus and the outer electrons.
• How does atomic radius affect ionisation energy?

  Electrostatic attraction will decrease as distance increases, so electrons in smaller atoms will be held closer to the nucleus, meaning the attraction is greater.
• How does electron shielding affect ionisation energy?

  Inner electron shells shield the outermost electrons from the full attractive force of the nucleus, reducing the effective nuclear charge experienced by the outer electrons.

  ->More electron shells provide more shielding.
• What increases ionisation energy?

  Low shielding and small atomic size, as the outermost electrons experience strong electrostatic attraction from the nucleus, and removing these tightly-held electrons requires a substantial energy input.
• What happens to ionisation energy down a group?

  Ionisation energy decreases down a group.
• Why does ionisation energy decrease down a group?

  • The atomic radius increases down the group, as more electron shells are added, so electrons are further from the nucleus.

  • The electron shielding increases down a group, as there are more inner electron shells to reduce nuclear attraction.
• Why does nuclear charge not have an affect on first ionisation energy decreasing down a group?

  • The nuclear charge increases down the group as more protons are added, increase attraction for electrons.

  • However, atomic radius and shielding effects down the group are greater than the nuclear charge effect.
• What happens to ionisation energy across periods?

  Ionisation energy increases across periods.
• Why does ionisation energy increase across periods?

  • The nuclear charge increases as more protons are added across periods.

  • The atomic radius decreases as extra electrons are only added to the same shell.

  • Electron shielding stays similar across periods with no extra inner shells.

  -> The increasing nuclear charge effect outweighs the similar shielding across periods, so ionisation energies generally increase across periods.
• What causes the drop in ionisation energy between groups 2 and 3?

  • In group 3, the electron is removed from a p orbital rather than an s orbital like in group 2.

  • p orbitals have slightly higher energy than s orbitals, so the outermost electron is on average further from the nucleus.

  • The p orbital also experiences addition shielding from the nucleus provided by the s electrons.

  
  

  -> As a result, less energy is required to remove the outermost p electron from the group 3 element compared to removing the outermost s electron from the group 2 element.
• What causes the drop in ionisation energy between groups 5 and 6?

  • In group 5, the electron is removed from a single occupied orbital.

  • In group 6, the electron is removed from an orbital containing two electrons.

  • The paired electrons in the group 6 element experience greater electron-electron repulsion.

  
  

  -> As a result, less energy is required to remove one of these paired electrons in the group 6 element compared to the unpaired electron in the group 5 element.
• How can electrons be removed from an atom?

  Electrons can be sequentially removed until only the bare nucleus remains.
• What is the successive ionisation energy?

  The energy required to remove each successive electron.
• What is the second ionisation energy?

  The energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions, to form 1 mole of gaseous 2+ ions.
• What is the equation for the second ionisation energy of lithium?

  Li+ (g) -> Li2+ (g) + e-
• What do successive ionisation energies provide evidence for?

  Electron shell structure
• What happens to successive ionisation energies within the same shell?

  Increase
• Why do successive ionisation energies increase within the same shell?

  • As successive electrons are removed from the same shell, the remaining electrons experience greater electrostatic attraction to the increasingly positive nucleus.

  • This increased nuclear attraction requires more energy to remove the next electron from that shell.
• What is an example of successive ionisation energies increasing within the same shell?

  For magnesium, the second ionisation energy (1.450 kJ mol-) is slightly higher than the first (740 kJ mol-), as these electrons are both being removed from the 3s subshell.